Chapter 16 Worksheet 1 (ws16.1)
Buffers and the Henderson-Hasselbalch Equation
1. State the operational and technical definitions of a buffer.
Operational definition (What does it do?):
Technical definition (How do you make one?):
2. The pH of a buffer is determined by two things:
3. Write equations that show what happens when a small amount of strong acid (H+) or strong base
(OH-) are added to a buffer.
4. Explain why the pH does not change very much when a small amount of strong acid or strong base
are added to a buffer.
5. What is meant by “buffering capacity”? What determines “buffering capacity”?
6. Most often (but not always) buffers are prepared such that the concentrations of the conjugate acid
and base are similar. Why?
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7. Calculate the pH of the following 3 solutions.
a. A 0.250 M solution of HF (Ka = 7.1 x 10-4)
b. A 0.500 M solution of NaF (Kb = ?)
c. A buffer that contains both 0.250 M HF and 0.500 M NaF
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8. The Henderson-Hasselbalch (H-H) equation provides a recipe for making a buffer of a given pH. It
is simply the logarithmic form of the Ka expression for the conjugate acid.
a. Derive the H-H equation below.
b. If [A-] = [HA] then pH ____________
c. If [A-] > [HA] then pH ____________
d. If [A-] < [HA] then pH ____________
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9. The Ka for HF is 7.1 x 10-4 so its pKa = ________________.
a. Mix 10.0 mL of 0.500 M HF with 10.0 mL of 0.500 M NaF. pH = ______________.
b. Mix 1.00 mL of 0.500 M HF with 10.0 mL of 0.500 M NaF. pH = ______________.
c. Mix 10.0 mL of 0.500 M HF mixed with 1.00 mL of 0.500 M NaF. pH = ______________.
d. A mixture of HF and NaF has a pH of 1.15.
[F  ]
 ____________.
[HF]
10. Use the Henderson-Hasselbalch equation to calculate the ratio of HCO3- to H2CO3 (Ka = 4.2 x 10-7)
in normal blood (which has a pH = 7.40).
11. Unlike the carbonate buffer used in blood, buffers used in a lab are usually prepared with a nearly
1:1 ratio of the acid and base. Circle the conjugate acid-base pair that you would choose to prepare a
buffer solution that has a pH of 4.50.
a. HClO and ClO(Ka = 3.5 x 10-8)
b. C6H5COOH and C6H5COO(Ka = 6.3 x 10-5)
c. HPO42- and PO43(Ka = 3.6 x 10-13)
Explain your choice:
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12. A buffer’s job is to prevent large pH changes upon the addition of small amounts of either strong
acid or strong base. As long as the buffer capacity is not exceeded, any added strong acid or strong
base will be neutralized by the buffer components. Let’s look at this process quantitatively.
a. What is the pH of a buffer made by mixing 10 mL of 0.40 M sodium dihyrdogen phosphate
with 10 mL of 0.40 M sodium hydrogen phosphate?
b. Write the net ionic equations for the neutralization reactions that occur upon addition of a
small amount of NaOH (or any metal hydroxide) or HCl (or any strong acid) to the buffer.
Add strong base:
Add strong acid:
c. What is the pH after 20.0 mL of 0.010 M NaOH or 0.010 M HCl is added to the buffer?
Remember that you can use moles rather than concentration in the H-H equation!
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